The table below shows the main oxidation states of the Group
elements. The major features of the redox chemistry of elements in this group are:
- Like oxygen, fluorine only has negative oxidation states.
- The most stable state for all elements is -1.
- All of the elements and the compounds containing halogen in positive oxidation states are oxidants.
- Despite being the most stable state for iodine, I– is a reasonable reductant in aqueous solution.
| | F | | Cl | | Br | | I |
*most stable state
| 0 (F2)
–1* (F–) | | +7 (ClO4–)
+5 (ClO3–)
+3 (ClO2–)
+1 (ClO–)
0 (Cl2)
–1* (Cl–) | | +7
+5
+1
0 (Br2)
–1* (Br–) | | +7
+5
+1
0 (I2)
–1* (I–) |
With the exception of fluorine, halogens and their oxoanions are susceptible to autooxidation-reduction (disproportionation).
Autooxidation-reduction (disproportionation) is a reaction in which atoms of the same element act as both the oxidant and the reductant. Therefore in order to undergo auto-oxidation-reduction reaction, an atom of an element must be in an intermediate oxidation state.
For example Cl2 (oxidation state 0) reacts with water to to form HOCl (Cl oxidation state +1) and HCl (Cl oxidation state -1).